Electrochemical table of metals. The world of modern materials - electrochemical series of voltages of metals

Sections: Chemistry, Competition "Presentation for the lesson"

Class: 11

Presentation for the lesson

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Targets and goals:

• Tutorial: Consideration of the chemical activity of metals based on the position in the periodic table D.I. Mendeleev and in the electrochemical voltage series of metals.
• Developing: Contribute to the development of auditory memory, the ability to compare information, think logically and explain ongoing chemical reactions.
• Educational: Forming a skill independent work, the ability to reasonably express one's opinion and listen to classmates, we instill in the children a sense of patriotism and pride in compatriots.

Equipment: PC with media projector, individual laboratories with a set of chemical reagents, models of crystal lattices of metals.

Lesson type: using technology for the development of critical thinking.

During the classes

I. Challenge stage.

Actualization of knowledge on the topic, the awakening of cognitive activity.

Bluff game: "Do you believe that ...". (Slide 3)

1. Metals occupy the upper left corner in the PSCE.
2. In crystals, metal atoms are bound by a metallic bond.
3. The valence electrons of metals are tightly bound to the nucleus.
4. Metals in the main subgroups (A) usually have 2 electrons in the outer level.
5. In the group from top to bottom there is an increase in the reducing properties of metals.
6. To assess the reactivity of a metal in solutions of acids and salts, it is enough to look at the electrochemical series of the voltages of metals.
7. To evaluate the reactivity of a metal in solutions of acids and salts, it is enough to look at the periodic table of D.I. Mendeleev

Question to the class? What does the entry mean? Me 0 - ne -\u003e Me + n(Slide 4)

Answer: Me0 - is a reducing agent, which means it interacts with oxidizing agents. The following can act as oxidizers:

1. Simple substances (+ O 2, Cl 2, S ...)
2. Complex substances (H 2 O, acids, salt solutions ...)

II. Understanding new information.

As a methodological technique, it is proposed to draw up a reference scheme.

Question to the class? What factors influence the reducing properties of metals? (Slide 5)

Answer: From the position in the periodic table of D.I. Mendeleev or from the position in the electrochemical series of the voltage of metals.

The teacher introduces the concepts: chemical activity and electrochemical activity.

Before starting the explanation, the children are invited to compare the activity of atoms To and Li position in the periodic table D.I. Mendeleev and the activity of simple substances formed by these elements according to their position in the electrochemical series of metal voltages. (Slide 6)

There is a contradiction:In accordance with the position of alkali metals in PSCE and according to the patterns of changes in the properties of elements in the subgroup, the activity of potassium is greater than that of lithium. In terms of position in the voltage series, lithium is the most active.

New material. The teacher explains the difference between chemical and electrochemical activity and explains that the electrochemical series of voltages reflects the ability of a metal to transform into a hydrated ion, where the measure of metal activity is energy, which consists of three terms (atomization energy, ionization energy and hydration energy). We write down the material in a notebook. (Slides 7-10)

Writing together in a notebook conclusion: The smaller the radius of the ion, the greater the electric field around it is created, the more energy is released during hydration, hence the stronger reducing properties of this metal in reactions.

History reference: presentation by a student on the creation by Beketov of a displacement series of metals. (Slide 11)

The action of the electrochemical voltage series of metals is limited only by the reactions of metals with electrolyte solutions (acids, salts).

Reminder:

1. The reducing properties of metals decrease during reactions in aqueous solutions under standard conditions (250°C, 1 atm.);
2. The metal to the left displaces the metal to the right of their salts in solution;
3. Metals standing up to hydrogen displace it from acids in solution (excl.: HNO3);
4. Me (to Al) + H 2 O -> alkali + H 2
   Other Me (up to H 2) + H 2 O -> oxide + H 2 (harsh conditions)
   Me (after H 2) + H 2 O -> do not react

(Slide 12)

The children are given notes.

Practical work:"Interaction of metals with salt solutions" (Slide 13)

Make the transition:

• CuSO4 —> FeSO4
• CuSO4 —> ZnSO4

Demonstration of the experience of interaction between copper and mercury (II) nitrate solution.

III. Reflection, contemplation.

We repeat: in which case we use the periodic table, and in which case a series of metal voltages is needed. (Slides 14-15).

We return to the initial questions of the lesson. We highlight questions 6 and 7 on the screen. We analyze which statement is not correct. On the screen - the key (check task 1). (Slide 16).

Summing up the lesson:

• What have you learned?
• In what case is it possible to use the electrochemical voltage series of metals?

Homework: (Slide 17)

1. To repeat the concept of "POTENTIAL" from the course of physics;
2. Finish the reaction equation, write the electronic balance equations: Cu + Hg (NO 3) 2 →
3. Given metals ( Fe, Mg, Pb, Cu)- offer experiments confirming the location of these metals in the electrochemical series of voltage.

We evaluate the results for the bluff game, work at the board, oral answers, communication, practical work.

Used Books:

1. O.S. Gabrielyan, G.G. Lysova, A.G. Vvedenskaya "Handbook for the teacher. Chemistry Grade 11, part II "Drofa Publishing House.
2. N.L. Glinka General Chemistry.

All metals, depending on their redox activity, are combined into a series called the electrochemical voltage series of metals (since the metals in it are arranged in order of increasing standard electrochemical potentials) or the activity series of metals:

Li, K, Ba, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, H 2 , Cu, Hg, Ag, Рt, Au

The most reactive metals are in the order of activity up to hydrogen, and the more to the left the metal is located, the more active it is. Metals that are next to hydrogen in the activity series are considered inactive.

Aluminum

Aluminum is a silvery white color. The main physical properties of aluminum are lightness, high thermal and electrical conductivity. In the free state, when exposed to air, aluminum is covered with a strong oxide film Al 2 O 3 , which makes it resistant to concentrated acids.

Aluminum belongs to the p-family metals. Electronic configuration of external energy level– 3s 2 3p 1 . In its compounds, aluminum exhibits an oxidation state equal to "+3".

Aluminum is obtained by electrolysis of the molten oxide of this element:

2Al 2 O 3 \u003d 4Al + 3O 2

However, due to the low yield of the product, the method of obtaining aluminum by electrolysis of a mixture of Na 3 and Al 2 O 3 is more often used. The reaction proceeds when heated to 960C and in the presence of catalysts - fluorides (AlF 3 , CaF 2 , etc.), while aluminum is released at the cathode, and oxygen is released at the anode.

Aluminum is able to interact with water after removing the oxide film from its surface (1), interact with simple substances (oxygen, halogens, nitrogen, sulfur, carbon) (2-6), acids (7) and bases (8):

2Al + 6H 2 O \u003d 2Al (OH) 3 + 3H 2 (1)

2Al + 3 / 2O 2 \u003d Al 2 O 3 (2)

2Al + 3Cl 2 = 2AlCl 3 (3)

2Al + N 2 = 2AlN (4)

2Al + 3S \u003d Al 2 S 3 (5)

4Al + 3C \u003d Al 4 C 3 (6)

2Al + 3H 2 SO 4 \u003d Al 2 (SO 4) 3 + 3H 2 (7)

2Al + 2NaOH + 3H 2 O \u003d 2Na + 3H 2 (8)

Calcium

In its free form, Ca is a silvery-white metal. When exposed to air, it instantly becomes covered with a yellowish film, which is the product of its interaction with the constituent parts of the air. Calcium is a fairly hard metal, has a cubic face-centered crystal lattice.

The electronic configuration of the external energy level is 4s 2 . In its compounds, calcium exhibits an oxidation state equal to "+2".

Calcium is obtained by electrolysis of molten salts, most often chlorides:

CaCl 2 \u003d Ca + Cl 2

Calcium is able to dissolve in water with the formation of hydroxides that exhibit strong basic properties (1), react with oxygen (2), forming oxides, interact with non-metals (3-8), dissolve in acids (9):

Ca + H 2 O \u003d Ca (OH) 2 + H 2 (1)

2Ca + O 2 \u003d 2CaO (2)

Ca + Br 2 \u003d CaBr 2 (3)

3Ca + N 2 \u003d Ca 3 N 2 (4)

2Ca + 2C = Ca 2 C 2 (5)

2Ca + 2P = Ca 3 P 2 (7)

Ca + H 2 \u003d CaH 2 (8)

Ca + 2HCl \u003d CaCl 2 + H 2 (9)

Iron and its compounds

Iron is a gray metal. In its pure form, it is quite soft, malleable and ductile. The electronic configuration of the external energy level is 3d 6 4s 2 . In its compounds, iron exhibits the oxidation states "+2" and "+3".

Metallic iron reacts with water vapor, forming a mixed oxide (II, III) Fe 3 O 4:

3Fe + 4H 2 O (v) ↔ Fe 3 O 4 + 4H 2

In air, iron is easily oxidized, especially in the presence of moisture (it rusts):

3Fe + 3O 2 + 6H 2 O \u003d 4Fe (OH) 3

Like other metals, iron reacts with simple substances, for example, halogens (1), dissolves in acids (2):

Fe + 2HCl \u003d FeCl 2 + H 2 (2)

Iron forms a whole range of compounds, since it exhibits several oxidation states: iron (II) hydroxide, iron (III) hydroxide, salts, oxides, etc. So, iron (II) hydroxide can be obtained by the action of alkali solutions on iron (II) salts without air access:

FeSO 4 + 2NaOH \u003d Fe (OH) 2 ↓ + Na 2 SO 4

Iron(II) hydroxide is soluble in acids and oxidized to iron(III) hydroxide in the presence of oxygen.

Salts of iron (II) exhibit the properties of reducing agents and are converted into iron (III) compounds.

Iron oxide (III) cannot be obtained by the combustion of iron in oxygen; to obtain it, it is necessary to burn iron sulfides or calcinate other iron salts:

4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2

2FeSO 4 \u003d Fe 2 O 3 + SO 2 + 3H 2 O

Iron (III) compounds exhibit weak oxidizing properties and are able to enter into OVR with strong reducing agents:

2FeCl 3 + H 2 S \u003d Fe (OH) 3 ↓ + 3NaCl

Iron and steel production

Steels and cast irons are alloys of iron with carbon, and the carbon content in steel is up to 2%, and in cast iron 2-4%. Steels and cast irons contain alloying additives: steels - Cr, V, Ni, and cast iron - Si.

There are various types of steels, so, according to their purpose, structural, stainless, tool, heat-resistant and cryogenic steels are distinguished. According to the chemical composition, carbon (low, medium and high carbon) and alloyed (low, medium and high alloyed) are distinguished. Depending on the structure, austenitic, ferritic, martensitic, pearlitic and bainitic steels are distinguished.

Steels have found application in many sectors of the national economy, such as construction, chemical, petrochemical, environmental protection, transport energy and other industries.

Depending on the form of carbon content in cast iron - cementite or graphite, as well as their quantity, several types of cast iron are distinguished: white (light color of the fracture due to the presence of carbon in the form of cementite), gray (gray color of the fracture due to the presence of carbon in the form of graphite). ), malleable and heat resistant. Cast irons are very brittle alloys.

The areas of application of cast iron are extensive - artistic decorations (fences, gates), body parts, plumbing equipment, household items (pans) are made from cast iron, it is used in the automotive industry.

Examples of problem solving

EXAMPLE 1

Exercise	An alloy of magnesium and aluminum weighing 26.31 g was dissolved in hydrochloric acid. In this case, 31.024 liters of colorless gas were released. Determine the mass fractions of metals in the alloy.
Solution	Both metals are capable of reacting with hydrochloric acid, as a result of which hydrogen is released: Mg + 2HCl \u003d MgCl 2 + H 2 2Al + 6HCl \u003d 2AlCl 3 + 3H 2 Find the total number of moles of hydrogen released: v(H 2) \u003d V (H 2) / V m v (H 2) \u003d 31.024 / 22.4 \u003d 1.385 mol Let the amount of substance Mg be x mol, and Al be y mol. Then, based on the reaction equations, we can write an expression for the total number of moles of hydrogen: x + 1.5y = 1.385 We express the mass of metals in the mixture: Then, the mass of the mixture will be expressed by the equation: 24x + 27y = 26.31 We got a system of equations: x + 1.5y = 1.385 24x + 27y = 26.31 Let's solve it: 33.24 -36y + 27y \u003d 26.31 v(Al) = 0.77 mol v(Mg) = 0.23mol Then, the mass of metals in the mixture: m (Mg) \u003d 24 × 0.23 \u003d 5.52 g m(Al) \u003d 27 × 0.77 \u003d 20.79 g Find the mass fractions of metals in the mixture: ώ =m(Me)/m sum ×100% ώ(Mg) = 5.52 / 26.31 × 100% = 20.98% ώ(Al) = 100 - 20.98 = 79.02%
Answer	Mass fractions of metals in the alloy: 20.98%, 79.02%

Attention! The description below is a reference material, it is not listed in this vinyl chart!

A SMALL COURSE OF ELECTROCHEMISTRY OF METALS

We have already become acquainted with the electrolysis of solutions of alkali metal chlorides and the production of metals using melts. Now let's try on a few simple experiments to study some of the laws of electrochemistry of aqueous solutions, galvanic cells, and also get acquainted with the production of protective galvanic coatings.
Electrochemical methods are used in modern analytical chemistry and serve to determine the most important quantities in theoretical chemistry.
Finally, the corrosion of metal objects, which causes great damage to the national economy, is in most cases an electrochemical process.

VOLTAGE RANGE OF METALS

The fundamental link for understanding electrochemical processes is the voltage series of metals. Metals can be arranged in a row that starts with reactive and ends with the least reactive noble metals:
Li, Rb, K, Ba, Sr, Ca, Mg, Al, Be, Mn, Zn, Cr, Ga, Fe, Cd, Tl, Co, Ni, Sn, Pb, H, Sb, Bi, As, Cu, Hg, Ag, Pd, Pt, Au.
This is how, according to the latest ideas, a series of voltages for the most important metals and hydrogen. If electrodes of a galvanic cell are made from any two metals of a row, then a negative voltage will appear on the material preceding in the row.
Voltage value ( electrochemical potential) depends on the position of the element in the voltage series and on the properties of the electrolyte.
We will establish the essence of the voltage series from a few simple experiments, for which we need a current source and electrical measuring instruments.

Metal coatings, "trees" and "ice patterns" without current

Let's dissolve about 10 g of crystalline copper sulfate in 100 ml of water and immerse a steel needle or a piece of iron sheet into the solution. (We recommend that you first clean the iron to a shine with a thin emery cloth.) After a short time, the iron will be covered with a reddish layer of released copper. The more active iron displaces the copper from the solution, with the iron dissolving as ions and the copper liberated as a metal. The process continues as long as the solution is in contact with the iron. As soon as the copper covers the entire surface of the iron, it will practically stop. In this case, a rather porous copper layer is formed, so that protective coatings cannot be obtained without the use of current.
In the following experiments, we will lower small strips of zinc and lead tin into the copper sulfate solution. After 15 minutes, take them out, rinse and examine under a microscope. We can see beautiful, ice-like patterns that are red in reflected light and consist of liberated copper. Here, too, more active metals transferred copper from the ionic to the metallic state.
In turn, copper can displace metals that are lower in the series of voltages, that is, less active. On a thin strip of sheet copper or flattened copper wire(having previously cleaned the surface to a shine), apply a few drops of silver nitrate solution. With the naked eye, it will be possible to notice the formed blackish coating, which under a microscope in reflected light looks like thin needles and plant patterns (the so-called dendrites).
To isolate zinc without current, it is necessary to use a more active metal. Excluding metals that violently interact with water, we find magnesium in the series of stresses above zinc. We place a few drops of zinc sulfate solution on a piece of magnesium tape or on a thin chip of an electron. zinc sulfate solutionwe get it by dissolving a piece of zinc in dilute sulfuric acid. Simultaneously with zinc sulfate, add a few drops of denatured alcohol. On magnesium, after a short period of time, we notice, especially under a microscope, zinc that has separated out in the form of thin crystals.
In general, any member of the voltage series can be forced out of solution, where it is in the form of an ion, and transferred to the metallic state. However, when trying all sorts of combinations, we may be disappointed. It would seem that if a strip of aluminum is immersed in solutions of salts of copper, iron, lead and zinc, these metals should stand out on it. But this, however, does not happen. The reason for the failure lies not in an error in the series of voltages, but is based on a special inhibition of the reaction, which in this case is due to a thin oxide film on the aluminum surface. In such solutions, aluminum is called passive.

LET'S LOOK BEYOND THE SCENE

In order to formulate the patterns of the ongoing processes, we can restrict ourselves to considering cations, and exclude anions, since they themselves do not participate in the reaction. (However, the type of anions affects the rate of deposition.) If, for simplicity, we assume that both the liberated and dissolved metals give doubly charged cations, then we can write:

Me 1 + Me 2 2+ = Me 1 2+ + Me 2

moreover, for the first experiment Me 1 = Fe, Me 2 = Сu.
So, the process consists in the exchange of charges (electrons) between atoms and ions of both metals. If we separately consider (as intermediate reactions) the dissolution of iron or the precipitation of copper, we get:

Fe = Fe 2+ + 2 e --

Сu 2+ + 2 e--=Cu

Now consider the case when the metal is immersed in water or in a salt solution, with the cation of which the exchange is impossible due to its position in the series of voltages. Despite this, the metal tends to go into solution in the form of an ion. In this case, the metal atom gives up two electrons (if the metal is divalent), the surface of the metal immersed in the solution is charged negatively with respect to the solution, and a double electric layer is formed at the interface. This potential difference prevents further dissolution of the metal, so that the process soon stops.
If two different metals are immersed in a solution, then they will both be charged, but the less active one is somewhat weaker, due to the fact that its atoms are less prone to splitting off electrons.
Connect both metals with a conductor. Due to the potential difference, the flow of electrons will flow from the more active metal to the less active one, which forms the positive pole of the element. A process takes place in which the more active metal goes into solution, and the cations from the solution are released on the more noble metal.

Essence of a galvanic cell

Let us now illustrate with a few experiments the above somewhat abstract reasoning (which, moreover, is a gross simplification).
First, fill a beaker with a capacity of 250 ml to the middle with a 10% sulfuric acid solution and immerse not too small pieces of zinc and copper into it. We solder or rivet a copper wire to both electrodes, the ends of which should not touch the solution.
As long as the ends of the wire are not connected to each other, we will observe the dissolution of zinc, which is accompanied by the release of hydrogen. Zinc, as follows from the voltage series, is more active than hydrogen, so the metal can displace hydrogen from the ionic state. Both metals form an electrical double layer. The potential difference between the electrodes is easiest to detect with a voltmeter. Immediately after turning on the device in the circuit, the arrow will indicate approximately 1 V, but then the voltage will quickly drop. If you connect a small light bulb to the element that consumes a voltage of 1 V, then it will light up - at first quite strongly, and then the glow will become weak.
By the polarity of the terminals of the device, we can conclude that the copper electrode is a positive pole. This can be proved even without a device by considering the electrochemistry of the process. Let us prepare a saturated solution of table salt in a small beaker or in a test tube, add about 0.5 ml of an alcohol solution of the phenolphthalein indicator and immerse both electrodes closed with a wire into the solution. Near the negative pole, a slight reddish coloration will be observed, which is caused by the formation of sodium hydroxide at the cathode.
In other experiments, one can place various pairs of metals in the cell and determine the resulting voltage. For example, magnesium and silver will give a particularly large potential difference due to the significant distance between them in a series of voltages, while zinc and iron, on the contrary, will give a very small one, less than a tenth of a volt. Using aluminum, we will not get practically any current due to passivation.
All these elements, or, as electrochemists say, circuits, have the disadvantage that when a current is taken, the voltage drops very quickly on them. Therefore, electrochemists always measure the true value of the voltage in a de-energized state using the voltage compensation method, that is, by comparing it with the voltage of another current source.
Let us consider the processes in the copper-zinc element in more detail. At the cathode, zinc goes into solution according to the following equation:

Zn = Zn2+ + 2 e --

Sulfuric acid hydrogen ions are discharged on the copper anode. They attach electrons coming through the wire from the zinc cathode and as a result, hydrogen bubbles are formed:

2H + + 2 e-- \u003d H 2

After a short period of time, copper will be covered with a thin layer of hydrogen bubbles. In this case, the copper electrode will turn into a hydrogen electrode, and the potential difference will decrease. This process is called electrode polarization. The polarization of the copper electrode can be eliminated by adding a little potassium dichromate solution to the cell after the voltage drop. After that, the voltage will increase again, since potassium dichromate will oxidize hydrogen to water. Potassium dichromate acts in this case as a depolarizer.
In practice, galvanic circuits are used, the electrodes of which are not polarized, or circuits, the polarization of which can be eliminated by adding depolarizers.
As an example of a non-polarizable element, consider the Daniell element, which was often used in the past as a current source. This is also a copper-zinc element, but both metals are immersed in different solutions. The zinc electrode is placed in a porous clay cell filled with dilute (about 20%) sulfuric acid. The clay cell is suspended in a large beaker containing a concentrated solution of copper sulfate, and at the bottom there is a layer of copper sulfate crystals. The second electrode in this vessel is a cylinder of copper sheet.
This element can be made from a glass jar, a commercially available clay cell (in extreme cases, use a flower pot, closing the hole in the bottom) and two electrodes of suitable size.
During the operation of the element, zinc dissolves with the formation of zinc sulfate, and copper ions are released on the copper electrode. But at the same time, the copper electrode is not polarized and the element gives a voltage of about 1 V. Actually, theoretically, the voltage at the terminals is 1.10 V, but when taking the current, we measure a slightly lower value, due to the electrical resistance of the cell.
If we do not remove the current from the cell, we must remove the zinc electrode from the sulfuric acid solution, because otherwise it will dissolve to form hydrogen.
A diagram of a simple cell, which does not require a porous partition, is shown in the figure. The zinc electrode is located in the glass jar at the top, and the copper electrode is located near the bottom. The entire cell is filled with a saturated sodium chloride solution. At the bottom of the jar we pour a handful of copper sulfate crystals. The resulting concentrated solution of copper sulfate will mix with the common salt solution very slowly. Therefore, during the operation of the cell, copper will be released on the copper electrode, and zinc in the form of sulfate or chloride will dissolve in the upper part of the cell.
Batteries now use almost exclusively dry cells, which are more convenient to use. Their ancestor is the Leclanchet element. The electrodes are a zinc cylinder and a carbon rod. The electrolyte is a paste that mainly consists of ammonium chloride. Zinc dissolves in the paste, and hydrogen is released on coal. To avoid polarization, the carbon rod is lowered into a linen bag with a mixture of coal powder and pyrolusite. The carbon powder increases the surface of the electrode, and the pyrolusite acts as a depolarizer, slowly oxidizing the hydrogen.
True, the depolarizing ability of pyrolusite is weaker than that of the previously mentioned potassium dichromate. Therefore, when current is received in dry cells, the voltage drops rapidly, they " get tired"due to polarization. Only after some time does the oxidation of hydrogen occur with pyrolusite. Thus, the elements" rest", if you do not pass current for some time. Let's check this on a flashlight battery, to which we connect a light bulb. Parallel to the lamp, that is, directly to the terminals, we connect a voltmeter.
At first, the voltage will be about 4.5 V. (Most often, three cells are connected in series in such batteries, each with a theoretical voltage of 1.48 V.) After a while, the voltage will drop, the light bulb will weaken. By reading the voltmeter, we can judge how long the battery needs to rest.
A special place is occupied by regenerating elements, known as accumulators. Reversible reactions take place in them, and they can be recharged after the cell is discharged by connecting to an external DC source.
Currently, lead-acid batteries are the most common; in them, the electrolyte is dilute sulfuric acid, into which two lead plates are immersed. The positive electrode is coated with lead dioxide PbO 2 , the negative electrode is metallic lead. The voltage at the terminals is approximately 2.1 V. During discharge, lead sulfate is formed on both plates, which again turns into metallic lead and into lead peroxide during charging.

PLATED COATINGS

The precipitation of metals from aqueous solutions with the help of an electric current is the reverse process of electrolytic dissolution, which we met when considering galvanic cells. First of all, let us examine the precipitation of copper, which is used in a copper coulometer to measure the amount of electricity.

Metal is deposited by current

Having bent the ends of two plates of thin sheet copper, we hang them on opposite walls of a beaker or, better, a small glass aquarium. We attach the wires to the plates with terminals.
Electrolyte prepare according to the following recipe: 125 g of crystalline copper sulfate, 50 g of concentrated sulfuric acid and 50 g of alcohol (denatured alcohol), the rest is water up to 1 liter. To do this, first dissolve copper sulfate in 500 ml of water, then carefully, in small portions, add sulfuric acid ( The heating! Liquid may splash!), then pour in alcohol and bring water to a volume of 1 liter.
We fill the coulometer with the prepared solution and include a variable resistance, an ammeter and a lead battery in the circuit. With the help of resistance, we adjust the current so that its density is 0.02-0.01 A/cm 2 of the electrode surface. If the copper plate has an area of ​​​​50 cm 2, then the current strength should be in the range of 0.5-1 A.
After some time, light red metallic copper will begin to precipitate at the cathode (negative electrode), and copper will go into solution at the anode (positive electrode). To clean the copper plates, we will pass a current in the coulometer for about half an hour. Then we take out the cathode, dry it carefully with filter paper and weigh it accurately. We install an electrode in the cell, close the circuit with a rheostat and maintain a constant current, for example 1 A. After an hour, we open the circuit and weigh the dried cathode again. At a current of 1 A per hour of operation, its mass will increase by 1.18 g.
Therefore, an amount of electricity equal to 1 ampere-hour, when passing through a solution, can release 1.18 g of copper. Or in general: the amount of substance released is directly proportional to the amount of electricity passed through the solution.
To isolate 1 equivalent of an ion, it is necessary to pass through the solution an amount of electricity equal to the product of the electrode charge e and the Avogadro number N A:
e*N A \u003d 1.6021 * 10 -19 * 6.0225 * 10 23 \u003d 9.65 * 10 4 A * s * mol -1 This value is indicated by the symbol F and is named after the discoverer of the quantitative laws of electrolysis Faraday number(exact value F- 96 498 A * s * mol -1). Therefore, to isolate a given number of equivalents from a solution n e through the solution, an amount of electricity equal to F*n e A * s * mol -1. In other words,
I*t =F*n e Here I- current, t is the time it takes for the current to pass through the solution. In chapter " Titration Basics"It has already been shown that the number of equivalents of a substance n e is equal to the product of the number of moles by the equivalent number:
n e = n*Z Consequently:

I*t = F*n*Z

In this case Z- ion charge (for Ag + Z= 1, for Cu 2+ Z= 2, for Al 3+ Z= 3, etc.). If we express the number of moles as the ratio of mass to molar mass ( n = m / M), then we get a formula that allows you to calculate all the processes that occur during electrolysis:

I*t =F*m*Z / M

Using this formula, you can calculate the current:

I = F*m*Z/(t*M)\u003d 9.65 * 10 4 * 1.18 * 2 / (3600 * 63.54) A * s * g * mol / (s * mol * g) \u003d 0.996 A

If we introduce the ratio for electrical work W email

W email = U*I*t and W email / U = I*t

then knowing the tension U, you can calculate:

W email = F*m*Z*U/M

You can also calculate how long it takes for the electrolytic release of a certain amount of a substance, or how much of a substance will be released in a certain time. During the experiment, the current density must be maintained within the specified limits. If it is less than 0.01 A / cm 2, then too little metal will be released, since copper (I) ions will be partially formed. If the current density is too high, the adhesion of the coating to the electrode will be weak, and when the electrode is removed from the solution, it may crumble.
In practice, galvanic coatings on metals are used primarily to protect against corrosion and to obtain a mirror finish.
In addition, metals, especially copper and lead, are refined by anodic dissolution and subsequent separation at the cathode (electrolytic refining).
To plate iron with copper or nickel, you must first thoroughly clean the surface of the object. To do this, polish it with elutriated chalk and sequentially degrease it with a dilute solution of caustic soda, water and alcohol. If the object is covered with rust, it is necessary to pickle it in advance in a 10-15% sulfuric acid solution.
We will hang the cleaned product in an electrolytic bath (a small aquarium or a beaker), where it will serve as a cathode.
The solution for applying copper plating contains 250 g of copper sulfate and 80-100 g of concentrated sulfuric acid in 1 liter of water (Caution!). In this case, a copper plate will serve as the anode. The surface of the anode should be approximately equal to the surface of the coated object. Therefore, you must always ensure that the copper anode hangs in the bath at the same depth as the cathode.
The process will be carried out at a voltage of 3-4 V (two batteries) and a current density of 0.02-0.4 A/cm 2 . The temperature of the solution in the bath should be 18-25 °C.
Pay attention to the fact that the plane of the anode and the surface to be coated are parallel to each other. It is better not to use objects of complex shape. By varying the duration of electrolysis, it is possible to obtain a copper coating of different thicknesses.
Preliminary copper plating is often resorted to in order to apply a durable coating of another metal to this layer. This is especially often used in iron chromium plating, zinc casting nickel plating and in other cases. True, very toxic cyanide electrolytes are used for this purpose.
To prepare an electrolyte for nickel plating, dissolve 25 g of crystalline nickel sulfate, 10 g of boric acid or 10 g of sodium citrate in 450 ml of water. Sodium citrate can be prepared by neutralizing a solution of 10 g citric acid diluted sodium hydroxide solution or soda solution. Let the anode be a nickel plate of the largest possible area, and take the battery as a voltage source.
The value of the current density with the help of a variable resistance will be maintained equal to 0.005 A/cm 2 . For example, with an object surface of 20 cm 2, it is necessary to work at a current strength of 0.1 A. After half an hour of work, the object will already be nickel plated. Take it out of the bath and wipe it with a cloth. However, it is better not to interrupt the nickel plating process, since then the nickel layer may passivate and the subsequent nickel coating will not adhere well.
In order to achieve a mirror shine without mechanical polishing, we introduce a so-called brightening additive into the plating bath. Such additives are, for example, glue, gelatin, sugar. You can enter into a nickel bath, for example, a few grams of sugar and study its effect.
To prepare an electrolyte for iron chromium plating (after preliminary copper plating), let's dissolve 40 g of CrO 3 chromic anhydride (Caution! Poison!) and exactly 0.5 g of sulfuric acid (in no case more!) in 100 ml of water. The process proceeds at a current density of about 0.1 A/cm 2 , and a lead plate is used as the anode, the area of ​​which should be slightly less than the area of ​​the chromium-plated surface.
Nickel and chrome baths are best heated slightly (up to about 35 °C). Please note that electrolytes for chromium plating, especially with a long process and high current strength, emit vapors containing chromic acid, which are very harmful to health. Therefore, chrome plating should be carried out under draft or outdoors, for example on a balcony.
In chromium plating (and, to a lesser extent, in nickel plating), not all of the current is used for metal deposition. At the same time, hydrogen is released. On the basis of a series of voltages, it would be expected that the metals standing in front of hydrogen should not be released from aqueous solutions at all, but, on the contrary, less active hydrogen should be released. However, here, as in the case of anodic dissolution of metals, the cathodic evolution of hydrogen is often inhibited and is observed only at high voltage. This phenomenon is called hydrogen overvoltage, and it is especially large, for example, on lead. Due to this circumstance, a lead battery can function. When the battery is charged, instead of PbO 2, hydrogen should appear on the cathode, but, due to overvoltage, hydrogen evolution begins when the battery is almost fully charged.

In chemistry textbooks, when presenting the topic "Acids", in one form or another, the so-called displacement series of metals is mentioned, the compilation of which is often attributed to Beketov.

For example, G. E. Rudzitis and F. G. Feldman, the once most widespread textbook for the 8th grade (from 1989 to 1995, it was published with a total circulation of 8.3 million copies), says the following. It is easy to verify from experience that magnesium reacts quickly with acids (using hydrochloric acid as an example), zinc reacts somewhat more slowly, iron even more slowly, and copper does not react with hydrochloric acid. “Similar experiments were carried out by the Russian scientist N. N. Beketov,” the authors of the textbook write further. – On the basis of experiments, he compiled a displacement series of metals: K, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb (H), Cu, Hg, Ag, Pt, Au. In this series, all metals that stand before hydrogen are able to displace it from acids. It is also reported that Beketov is “the founder of physical chemistry. In 1863 he compiled a displacement series of metals, which is named after the scientist. Next, students are told that in the Beketov series, metals to the left displace metals to the right from solutions of their salts. The exception is the most active metals. Similar information can be found in other school textbooks and manuals, for example: “The Russian chemist N. N. Beketov investigated all metals and arranged them according to their chemical activity in a displacement series (activity series)”, etc.

Several questions may arise here.

Question one. Didn't chemists know before Beketov's experiments (that is, before 1863) that magnesium, zinc, iron, and a number of other metals react with acids to release hydrogen, while copper, mercury, silver, platinum, and gold do not possess this property?

Question two. Didn't chemists before Beketov notice that some metals can displace others from solutions of their salts?

Question three. In the book by V. A. Volkov, E. V. Vonsky, G. I. Kuznetsov “Outstanding chemists of the world. Biographical Reference Book (Moscow: Vysshaya Shkola, 1991) says that Nikolai Nikolaevich Beketov (1827–1911) is “a Russian physical chemist, academician… one of the founders of physical chemistry… He studied the behavior of organic acids at high temperatures. Synthesized (1852) benzureide and aceturide. Put forward (1865) a number of theoretical provisions on the dependence of the direction of reactions on the state of the reagents and external conditions ... Determined the heat of formation of oxides and chlorides of alkali metals, for the first time received (1870) anhydrous oxides of alkali metals. Using the ability of aluminum to restore metals from their oxides, he laid the foundations of aluminothermy ... President of the Russian Physico-Chemical Society .... ". And not a word about his compilation of a displacement series, which was included (unlike, for example, ureides - urea derivatives) in school textbooks published in millions of copies!

It is hardly necessary to blame the authors of the biographical guide for forgetting the important discovery of the Russian scientist: after all, D. I. Mendeleev, who by no means can be reproached for unpatriotism, in his classic textbook "Fundamentals of Chemistry" also never mentions Beketov's displacement series, although 15 times refers to various of his works. To answer all these questions, we will have to make an excursion into the history of chemistry, to figure out who and when proposed the activity series of metals, what experiments N. N. Beketov himself conducted and what his displacement series is.

The first two questions can be answered in the following way. Of course, both the release of hydrogen from acids by metals, and various examples their displacement of each other from the salts were known long before the birth of Beketov. For example, in one of the manuals of the Swedish chemist and mineralogist Thornburn Olaf Bergman, published in 1783, it is recommended to displace lead and silver from solutions using iron plates when analyzing polymetallic ores. When carrying out calculations on the iron content in the ore, one should take into account that part of it that passed into the solution from the plates. In the same manual, Bergman writes: “Metals can be displaced from solutions of their salts by other metals, and some consistency is observed. In the series of zinc, iron, lead, tin, copper, silver and mercury, zinc displaces iron, etc.” And, of course, it was not Bergman who first discovered these reactions: such observations date back to alchemical times. The most famous example of such a reaction was used in the Middle Ages by charlatans who publicly demonstrated the "transformation" of an iron nail into red "gold" when they dipped the nail into a solution of copper sulphate. Now this reaction is demonstrated in chemistry classes at school. What is the essence of Beketov's new theory? Before the advent of chemical thermodynamics, chemists explained the flow of a reaction in one direction or another by the concept of the affinity of some bodies for others. The same Bergman, based on well-known displacement reactions, developed from 1775 the theory of selective affinity. According to this theory, the chemical affinity between two substances under given conditions remains constant and does not depend on the relative masses of the reactants. That is, if bodies A and B are in contact with body C, then the body that has a greater affinity for it will connect with C. For example, iron has a greater affinity for oxygen than mercury, and therefore it will be the first to be oxidized by it. It was assumed that the direction of the reaction is determined solely by the chemical affinity of the reacting bodies, and the reaction goes to the end. Bergman compiled tables of chemical affinity, which were used by chemists until the beginning of the 19th century. These tables included, in particular, various acids and bases.

Almost simultaneously with Bergman, the French chemist Claude Louis Berthollet developed another theory. Chemical affinity was also associated with the attraction of bodies to each other, but other conclusions were drawn. By analogy with the law of universal attraction, Berthollet believed that in chemistry, attraction should also depend on the mass of the reacting bodies. Therefore, the course of the reaction and its result depend not only on the chemical affinity of the reagents, but also on their quantities. For example, if bodies A and B can react with C, then body C will be distributed between A and B according to their affinities and masses, and not a single reaction will reach the end, since equilibrium will come when AC, BC and free A and B coexist simultaneously. It is very important that the distribution of C between A and B can vary depending on the excess of A or B. Therefore, with a large excess, a body with low affinity can almost completely “select” body C from its “rival”. But if one of the reaction products (AC or BC) is removed, then the reaction will go to the end and only the product that leaves the scope is formed.

Berthollet made his conclusions by observing the processes of precipitation from solutions. These conclusions sound surprisingly modern, apart from outdated terminology. However, Berthollet's theory was qualitative; it did not provide a way to measure affinity values.

Further advances in theory were based on discoveries in the field of electricity. Italian physicist Alessandro Volta late XVIII in. showed that when different metals come into contact, an electric charge arises. Conducting experiments with various pairs of metals and determining the sign and magnitude of the charge of some metals in relation to others, Volta established a series of voltages: Zn, Pb, Sn, Fe, Cu, Ag, Au. Using pairs of different metals, Volta designed a galvanic cell, the strength of which was the greater, the farther apart the members of this series were. The reason for this was unknown at the time. True, back in 1797, the German scientist Johann Wilhelm Ritter predicted that metals should be in the series of stresses in order of decreasing their ability to combine with oxygen. In the case of zinc and gold, this conclusion was not in doubt; as for other metals, it should be noted that their purity was not very high, so the Volta series does not always correspond to the modern one.

Theoretical views on the nature of the processes occurring in this case were very vague and often contradictory. The famous Swedish chemist Jöns Jakob Berzelius at the beginning of the 19th century. created an electrochemical (or dualistic, from lat. dualis - "dual") the theory of chemical compounds. In accordance with this theory, it was assumed that each chemical compound consists of two parts - positively and negatively charged. In 1811, Berzelius, based on the chemical properties of the elements known to him, arranged them in a row so that each term in it was electronegative with respect to the previous one and electropositive with respect to the next. In an abbreviated version, the following were assigned to the electronegative elements (in descending order):

O, S, N, Cl, Br, S, Se P, As, Cr, B, C, Sb, Te, Si.

Then followed the transition element - hydrogen, and after it - electropositive elements (in order of increasing this property):

Au, Pt, Hg, Ag, Cu, Bi, Sn, Pb, Cd, Co, Ni, Fe, Zn, Mn, Al, Mg, Ca, Sr, Ba, Li, Na, K.

This series, if you rewrite all the metals in reverse order, is very close to the modern one. Some differences in the order of the metals in this series are probably due to the insufficient purification of substances in the time of Berzelius, as well as some other properties of the metals that Berzelius was guided by. According to Berzelius, the farther the elements are from each other in this series, the more opposite electric charges they have and the more durable they form chemical compounds with each other.

Berzelius' theory of dualism in the middle of the 19th century. was dominant. Its failure was shown by the founders of thermochemistry, the French scientist Marcellin Berthelot and the Danish researcher Julius Thomsen. They measured chemical affinity by the work that a chemical reaction can produce. In practice, it was measured by the heat of the reaction. These works led to the creation of chemical thermodynamics, a science that made it possible, in particular, to calculate the position of equilibrium in a reacting system, including equilibrium in electrochemical processes. The theoretical basis for the activity series (and for the stress series) in solutions was laid at the end of the 19th century. German physical chemist Walter Nernst. Instead of a qualitative characteristic - the affinity or ability of a metal and its ion to certain reactions - an exact quantitative value appeared that characterizes the ability of each metal to pass into solution in the form of ions, and also to be reduced from ions to metal on the electrode. Such a value is the standard electrode potential of the metal, and the corresponding series, arranged in order of potential changes, is called the series of standard electrode potentials. (The standard state assumes that the concentration of ions in the solution is 1 mol/l, and the gas pressure is 1 atm; most often, the standard state is calculated for a temperature of 25 ° C.)

The standard potentials of the most active alkali metals were calculated theoretically, since it is impossible to measure them experimentally in aqueous solutions. To calculate the potentials of metals at different concentrations of their ions (i.e., in non-standard states), the Nernst equation is used. Electrode potentials have been determined not only for metals, but also for many redox reactions involving both cations and anions. This makes it possible to theoretically predict the possibility of a variety of redox reactions occurring under various conditions. It should also be noted that in non-aqueous solutions, the potentials of the metals will be different, so that the sequence of metals in the series may change markedly. For example, in aqueous solutions, the potential of the copper electrode is positive (+0.24 V) and copper is located to the right of hydrogen. In a solution of acetonitrile CH3CN, the copper potential is negative (–0.28 V), i.e., copper is located to the left of hydrogen. Therefore, the following reaction takes place in this solvent: Cu + 2HCl = CuCl2 + H2.

Now it's time to answer the third question and find out what exactly Beketov studied and what conclusions he came to.

One of the most prominent Russian chemists, N. N. Beketov, after graduating (in 1848) from Kazan University, worked for some time at the Medical and Surgical Academy in the laboratory of N. N. Vinin, then at St. Kharkov University. Shortly after receiving the university department of chemistry in 1857, Beketov went abroad for a year “with an appointment of a thousand rubles a year in excess of the salary received” - at that time it was a large amount. During his stay in Paris, he published (in French) the results of his earlier studies in Russia on the displacement of certain metals from solutions by hydrogen and on the reducing effect of zinc vapor. At a meeting of the Paris Chemical Society, Beketov reported on his work on the reduction of SiCl4 and BF3 with hydrogen. These were the first links in the chain of research devoted to the displacement of some elements by others, which Beketov began in 1856 and completed in 1865.

Already abroad, Beketov drew attention to himself. It is enough to quote the words of D. I. Mendeleev, whom Beketov met in Germany: “From Russian chemists abroad, I learned Beketov ... Savich, Sechenov. That's all ... people who do honor to Russia, people with whom I am glad that I got along.

In 1865, Beketov's dissertation "Research on the phenomena of displacement of some elements by others" was published in Kharkov. This work was republished in Kharkov in 1904 (in the collection “In memory of the 50th anniversary of the scientific activity of N. N. Beketov”) and in 1955 (in the collection “N. N. Beketov. Selected Works in Physical Chemistry”) .

Let's get acquainted with this work of Beketov in more detail. It consists of two parts. The first part (it contains six sections) presents the results of the author's experiments in great detail. The first three sections are devoted to the action of hydrogen on solutions of silver and mercury salts at various pressures. It seemed to Beketov an extremely important task to find out the place of hydrogen in a series of metals, as well as the dependence of the direction of the reaction on external conditions - pressure, temperature, concentration of reagents. He conducted experiments both in solutions and with dry substances. It was well known to chemists that hydrogen easily displaces some metals from their oxides at high temperatures, but is inactive at low temperatures. Beketov found that the activity of hydrogen increases with increasing pressure, which he associated with the "greater density" of the reagent (now they would say - with a higher pressure, i.e., gas concentration).

Studying the possibility of displacing metals with hydrogen from solutions, Beketov set up a number of rather risky experiments. For the first time in the history of chemistry, Beketov applied pressures exceeding 100 atm. He conducted experiments in the dark, in sealed glass tubes with several bends (elbows). In one knee he placed a solution of salt, in the other - acid, and at the end of the tube - metallic zinc. By tilting the tube, Beketov made the zinc fall into the acid taken in excess. Knowing the mass of dissolved zinc and the volume of the tube, it was possible to estimate the achieved hydrogen pressure. In some experiments, Beketov specified the pressure by the degree of compression of air by a liquid in a thin capillary soldered to a tube. The opening of the tube was always accompanied by an explosion. In one of the experiments, in which the pressure reached 110 atm, an explosion during the opening of the tube (it was carried out in water under an overturned cylinder) shattered a thick-walled cylinder, the volume of which was a thousand times greater than the volume of the tube with reagents.

Experiments have shown that the action of hydrogen depends not only on its pressure, but also on the "strength of the metallic solution", that is, on its concentration. The reduction of silver from the ammonia solution of AgCl begins even before the complete dissolution of zinc at a pressure of about 10 atm - the transparent solution turns brown (first at the border with the gas, then throughout the mass), and after a few days gray silver powder settles on the walls. No reaction was observed at atmospheric pressure. Silver was also reduced from nitrate and sulfate, and hydrogen acted on silver acetate at atmospheric pressure. Metal balls were released from mercury salts at high pressure, but copper and lead nitrates could not be reduced even at high hydrogen pressure. The reduction of copper was observed only in the presence of silver and platinum at pressures up to 100 atm. Beketov used platinum to speed up the process, that is, as a catalyst. He wrote that platinum is more conducive to the displacement of certain metals than pressure, since hydrogen on the surface of platinum "is subject to greater attraction and should have the greatest density." Now we know that hydrogen adsorbed on platinum is activated due to its chemical interaction with metal atoms.

In the fourth section of the first part, Beketov describes experiments with carbon dioxide. He studied its effect on solutions of calcium acetate at different pressures; discovered that the reverse reaction - the dissolution of marble in acetic acid at a certain gas pressure stops even with an excess of acid.

In the last sections of the experimental part, Beketov described the effect of zinc vapor at high temperatures on compounds of barium, silicon, and aluminum (he calls the latter element clay, as was customary in those years). By reducing silicon tetrachloride with zinc, Beketov was the first to obtain sufficiently pure crystalline silicon. He also found that magnesium reduces aluminum from cryolite (sodium fluoroaluminate "in-house") and silicon from its dioxide. In these experiments, the ability of aluminum to restore barium from oxide and potassium from hydroxide was also established. So, after calcining aluminum with anhydrous barium oxide (with a small addition of barium chloride to lower the melting point), an alloy was formed, which, according to the results of the analysis, is 33.3% barium, the rest is aluminum. At the same time, many hours of calcining aluminum with powdered barium chloride did not lead to any changes.

The unusual reaction of aluminum with KOH was carried out in a curved gun barrel, in the closed end of which pieces of KOH and aluminum were placed. With a strong incandescence of this end, potassium vapor appeared, which condensed in the cold part of the barrel, "from where a few pieces of soft metal were obtained, burning with a violet flame." Rubidium and cesium were later isolated in a similar way.

The second part of Beketov's work is devoted to the theory of the displacement of some elements by others. In this part, Beketov first analyzed numerous experimental data - both his own and those conducted by other researchers, including Breslav Professor Fischer, as well as Davy, Gay-Lussac, Berzelius, Wöhler. Of particular note are "several interesting facts about the precipitation of metals by the wet route" discovered by the English chemist William Odling. At the same time, Beketov considers the cases of displacement of some elements by others "wet way", i.e. in solutions, and "dry way", i.e. during calcination of reagents. This was logical, since it is impossible to experimentally carry out reactions in aqueous solutions involving alkali and alkaline earth metals, since they actively react with water.

Then Beketov sets out his theory, designed to explain the different activity of the elements. Having arranged all the metals in a row according to their specific gravity (i.e., density), Beketov found that it agrees quite well with the known displacement series. “Consequently,” concludes Beketov, “the place of the metal ... in the displacement series can be fairly correctly determined and, so to speak, predicted in advance by its specific gravity.” Some uncertainty is observed only between "metals adjacent in specific gravity". Thus, potassium is usually a "more energetic" element and, for example, displaces sodium from NaCl when calcined, although potassium is more volatile. However, reverse processes are also known: for example, sodium can displace potassium from its hydroxide and acetate. “As for the ratio of the first alkaline group to the second and the ratio of the metals of the second group to each other, they are still little studied,” writes Beketov.

Beketov met with more serious difficulties. For example, he succeeded in reducing zinc with aluminum from a ZnCl2 solution and failed from a ZnSO4 solution. In addition, aluminum "absolutely did not restore iron, nickel, cobalt, cadmium from solutions." Beketov explained this by the fact that aluminum "acts mainly on water", and assumed that these reactions should go in the absence of water - "dry way". Indeed, later Beketov discovered such reactions and actually discovered aluminothermy.

Another difficulty was that some metals fell out of the rule of specific gravity. So, copper (density 8.9) in the activity series is located not before, but after lead (density 11.4 - Beketov's density values ​​are slightly different from modern ones). Such an "anomaly" forced Beketov to try to displace the more active lead with less active copper. He placed copper plates in hot saturated solutions of lead chloride - neutral and acidic, in an ammonia solution of lead oxide, heated copper with dry oxide and lead chloride. All experiments were unsuccessful, and Beketov was forced to admit "retreat from general rule". Other "anomalies" concerned silver (density 10.5) and lead, as well as silver and mercury (density 13.5), since both lead and mercury reduce the "lighter" silver from solutions of its salts. Beketov explained the anomaly with mercury by the fact that this metal is liquid and therefore its activity is higher than follows from the rule of specific gravity.

Beketov extended his rule to non-metals. For example, in the series chlorine (density of liquid chlorine 1.33), bromine (density 2.86), iodine (density 4.54), the lightest element is at the same time the most active (fluorine was obtained by Moissan only 20 years later). The same is observed in the series O, S, Se, Te: oxygen is the most active and quite easily displaces the rest of the elements from their compounds with hydrogen or with an alkali metal.

Beketov explained his rule by analogy with mechanics: the specific gravity is related to the mass of particles (ie, atoms) and the distance between them in a simple substance. Knowing the densities of metals and their relative atomic masses, one can calculate the relative distances between atoms. The greater the distance between them, the easier, according to Beketov, the atoms are separated in chemical processes. This is also connected with the mutual "affinity" of various elements, and the ability to displace each other from compounds. Having calculated the relative distance between atoms in different metals and taking potassium as a standard, Beketov obtained the following values: K - 100, Na - 80, Ca - 65, Mg - 53, Al - 43, etc. up to platinum.

Further summary Beketov’s theory regarding the relative strength of chemical compounds (namely, the ability of some elements to displace others is connected with this) can be found in D. I. Mendeleev’s textbook “Fundamentals of Chemistry” (quoted from the 1947 edition using modern terminology): “... Professor N. N. Beketov, in his work “Investigations on the Phenomena of Repression” (Kharkov, 1865), proposed a special hypothesis, which we will state almost in the words of the author.

For aluminum oxide Al2O3 is stronger than halides AlCl3 and AlI3. In the oxide, the ratio Al: O = 112: 100, for chloride Al: Cl = 25: 100, for iodide Al: I = 7: 100. For silver oxide Ag2O (ratio 1350: 100) is less durable than chloride (Ag: Cl = = 100: 33), and iodide is the most durable (Ag: I = 85: 100). From these and similar examples it can be seen that the most durable are those compounds in which the masses of the connecting elements become almost the same. Therefore, there is a desire of large masses to combine with large ones, and small masses - with small ones, for example: Ag2O + 2KI give K2O + 2AgI. For the same reason when elevated temperatures Ag2O, HgO, Au2O3 and similar oxides, composed of unequal masses, decompose, while oxides of light metals, as well as water, do not decompose so easily. The most heat-resistant oxides - MgO, CaO, SiO2, Al2O3 approach the mass equality condition. For the same reason, HI decomposes more easily than HCl. Chlorine does not act on MgO and Al2O3, but acts on CaO, Ag2O, etc.

To understand the true relationships of affinities, - concludes Mendeleev, - those additions to the mechanical theory of chemical phenomena that Beketov gives are still far from enough. Nevertheless, in his way of explaining the relative strength of many compounds, one can see a very interesting statement of questions of paramount importance. Without such attempts it is impossible to grasp the complex objects of experiential knowledge.

So, without belittling the merits of the remarkable chemist, it should be recognized that, although the theory of N. N. Beketov played a significant role in the development of theoretical chemistry, one should not attribute to him the establishment of the relative activity of metals in the reaction of displacement of hydrogen from acids and the corresponding series of activity of metals: its mechanical The theory of chemical phenomena remained in the history of chemistry as one of its many stages.

Why, then, in some books, Beketov is credited with something that he did not discover? This tradition, like many others, probably appeared in the late 1940s and early 1950s. of the 20th century, when a campaign to combat “complaining to the West” was raging in the USSR, and the authors simply had to attribute all more or less noticeable discoveries in science exclusively to domestic scientists, and even citing foreign authors was considered sedition (it was in those years that the joke about that “Russia is the birthplace of elephants”). For example, M. V. Lomonosov was credited with the discovery of the law of conservation of energy, which was discovered only in the middle of the 19th century. Here is a concrete example of the presentation of the history of science of those times. In Vladimir Orlov’s book “On a Courageous Thought” (M.: Molodaya Gvardiya, 1953), inventions in the field of electricity are described in the following words: “Foreigners ruined the cradle of electric light ... The Americans stole a wonderful Russian invention ... Edison in America greedily began to improve the Russian invention ... Foreign scientists cripple an electric lamp created by the genius of the Russian people ... The American imperialists disgraced electricity ... Following them, the Yugoslav fascists disgraced electric light ... "- etc., etc. Separate echoes of those bad memory of the times, apparently, remained in some textbooks, and they should be disposed of. As one of the historians of chemistry said, "Lomonosov is great enough not to attribute other people's discoveries to him."

"The candle was burning..."

The phenomena observed during the burning of a candle are such that there is not a single law of nature that would not be affected in one way or another.

Michael Faraday. History of the candle

This story is about "experimental investigation". The main thing in chemistry is experiment. In laboratories around the world, millions of various experiments have been and continue to be performed, but it is extremely rare for a professional researcher to do it the way some young chemists do: what if something interesting happens? Most often, the researcher has a clearly formulated hypothesis, which he seeks to either confirm or disprove experimentally. But now the experience is over, the result is obtained. If it does not agree with the hypothesis, then it is incorrect (of course, if the experiment is set up correctly and it is reproduced several times). What if it agrees? Does this mean that the hypothesis is correct and it is time to transfer it into the category of theory? A novice researcher sometimes thinks so, but an experienced one does not rush to conclusions, but first thinks firmly whether it is possible to explain the result obtained in some other way.

The history of chemistry knows thousands of examples of how such "thinking" is useful. The next three stories are just devoted to how dangerous it can be to believe that a "successful" experiment proves the correctness of the hypothesis. Sometimes in the classroom they show such an experience. A small wooden or foam circle is allowed to float in a plate of water, on which a burning candle is fixed. An inverted glass jar is lowered onto a circle with a candle and placed in this position on the bottom of the plate. After a while, the candle goes out, and part of the jar is filled with water. This experiment is supposed to show that only a fifth of the air (oxygen) supports combustion. Indeed, at first glance it looks like the water has risen by about a fifth, although more accurate measurements are not usually taken. At first glance, the experiment is simple and quite convincing: after all, oxygen in the air is indeed 21% by volume. However, from the point of view of chemistry, it is not all right. Indeed, candles are made from paraffin, and paraffin consists of saturated hydrocarbons of composition C n H2 n+2 with 18–35 carbons. The combustion reaction equation can be written in general form as follows: n H2 n +2 + (3 n+ 1)/2 O2 → n CO2 + ( n+ 1)H2O. Because n is large, then the coefficient in front of oxygen is very close to 1.5 n(for n= 18 difference between (3 n+ +1)/2 and 1.5 n will be less than 2%, for n= 30 it will be even less). Thus, for 1.5 volume of oxygen consumed, 1 volume of CO2 is released. Therefore, even if all the oxygen from the can (it is 0.21 by volume there) is used up, then instead of it, after combustion, 0.21: 1.5 = 0.14 volumes of carbon dioxide should be released. This means that water should not fill a fifth of the jar at all!

But is this reasoning correct? After all, carbon dioxide, as you know, is highly soluble in water. Maybe it will all “go into the water”? However, the process of dissolving this gas is very slow. This was shown by special experiments: pure water in an inverted jar filled with CO2 almost does not rise in an hour. The experiment with the candle lasts less than a minute, therefore, even if oxygen is completely used up, water should enter the jar by only 0.21 - 0.1 = 0.07 of its volume (about 7%).

But that's not all. It turns out that the candle “burns” in the jar not all the oxygen, but only a small part of it. An analysis of the air in which the candle went out showed that it still contained 16% oxygen (interestingly, the oxygen content in a normal human exhalation decreases to about the same level). This means that water should not enter the jar at all! Experience, however, shows that this is not the case. How to explain it?

The simplest assumption: a burning candle heats up the air, its volume increases, and part of the air comes out of the jar. After cooling the air in the jar (this happens quite quickly), the pressure in it decreases, and water enters the jar under the action of external atmospheric pressure. In accordance with the law of ideal gases (and air in the first approximation can be considered an ideal gas), in order for the volume of air to increase by 1/5, its temperature (absolute) must also increase by 1/5, i.e. increase from 293 K (20 ° C) up to 1.2 293 = 352 K (about 80 ° C). Not so much! Heating the air with a candle flame at 60° is quite possible. It remains only to check experimentally whether air comes out of the jar during the experiment.

The first experiments, however, did not seem to confirm this assumption. So, in a series of experiments carried out with a wide-mouth jar with a volume of 0.45 l, there were no signs of “gurgling” of air from under the edge of the jar. Another unexpected observation: the water in the jar, while the candle was burning, almost did not enter.

And only after the candle went out, the water level in the inverted jar quickly rose. How to explain it?

It could be assumed that while the candle is burning, the air in the jar heats up, but at the same time, not its volume increases, but the pressure, which prevents the water from being sucked in. After the combustion stops, the air in the jar cools down, its pressure drops, and the water rises. However, this explanation does not fit. First, water is not heavy mercury, which would keep air from escaping a jar with a slight increase in pressure. (The mercury seal was once used by all physicists and chemists who studied gases.) Indeed, water is 13.6 times lighter than mercury, and the height of the water seal between the edge of the jar and the level of water in the plate is small. Therefore, even a small increase in pressure would inevitably cause air to "bubbling" through the valve.

The second objection is even more serious. Even if the water level in the plate were higher and the water would not release heated air under high pressure from the jar, then after the air in the jar cools down, both its temperature and pressure would return to their original values. So there would be no reason for air to enter the jar.

The riddle was solved only by changing a small detail during the experiment. Usually the jar is “put on” on top of the candle. So, maybe this is the reason for the strange behavior of the air in the bank? A burning candle creates an upward flow of heated air, and as the jar moves from above, the hot air displaces colder air from the jar before the edge of the jar touches the water. After that, the air temperature in the jar, while the candle is burning, changes little, so the air does not leave it (and also does not go inside). And after the cessation of combustion and cooling of the hot air in the jar, the pressure in it noticeably decreases, and the external atmospheric pressure drives part of the water into the jar.

To test this assumption, in several experiments, the jar was “put on” on the candle not from above, but from the side, almost touching the flame with the edge of the jar, after which, with a quick downward movement, the jar was placed on the bottom of the plate. And immediately from under the edge of the jar, air bubbles began to rapidly emerge! Naturally, after the burning of the candle stopped, the water was sucked inward - approximately to the same level as in previous experiments.

So this experiment with a candle cannot in any way illustrate the composition of air. But he reiterates wise saying great physicist, rendered in the epigraph.

Getting closer to balance...

Let us consider one more erroneous explanation of the experiment, in which gases are also heated. This explanation has found its way into popular chemistry articles and even college textbooks. So, in a number of foreign textbooks on general chemistry, a beautiful experiment is described, the essence of which we will illustrate with a quote from Noel Waite's textbook "Chemical Kinetics". Relaxation method. The Eigen method, for which the author was awarded in 1967 Nobel Prize in chemistry, is called the relaxation method. The reacting system reaches a state of equilibrium under certain conditions. These conditions (temperature, pressure, electric field) are then quickly violated - faster than the equilibrium is shifted. The system again comes into equilibrium, but now under new conditions; this is called "relaxing to a new equilibrium position". While relaxation is taking place, a change in some property of the system is monitored...

An experiment demonstrating the phenomenon of relaxation.

In some cases, the state of equilibrium is established so slowly under new conditions that the change in concentration can be followed with the help of conventional laboratory equipment and thus the phenomenon of relaxation can be observed. As an example, consider the transition of nitrogen dioxide (dark brown gas) to a dimer (colorless gas):

Fill the glass gas syringe with approximately 80 cm3 of gas. Quickly press the plunger of the syringe and compress the gas to 50–60 cm3. Verify that the color of the gas has changed. First there will be a rapid darkening of the gas, as the concentration of NO2 will increase, but then a slow brightening will occur, since high pressure contributes to the formation of N2O4, and equilibrium will be reached under new external conditions.

In a number of textbooks, a similar description is given to illustrate Le Chatelier's principle: with increasing gas pressure, the equilibrium shifts towards a decrease in the number of molecules, in this case towards a colorless N2O4 dimer. The text is accompanied by three color photographs. They show how, immediately after compression, the initially yellowish-brown mixture becomes dark brown, and in the third photograph, taken after a few minutes, the gas mixture in the syringe noticeably brightens.

Sometimes they add that the piston must be pressed as quickly as possible so that the balance does not have time to move during this time.

At first glance, this explanation looks very convincing. However, a quantitative examination of the processes in the syringe completely refutes all conclusions. The fact is that the indicated equilibrium between nitrogen dioxide NO2 and its dimer (nitrogen tetroxide) N2O4 is established extremely quickly: in millionths of a second! Therefore, it is impossible to compress the gas in the syringe faster than this equilibrium is established. Even if you move the piston in the steel "syringe" with the help of an explosion, the equilibrium would most likely have time to be established as the piston moves due to its inertia. How else can the phenomenon observed in this experiment be explained? Of course, a decrease in volume and a corresponding increase in the concentration of gases leads to an increase in color. But not this main reason. Anyone who has inflated a bicycle tube with a hand pump knows that a pump (especially an aluminum one) gets very hot. The friction of the piston on the pump tube has nothing to do with it - this is easy to verify by making a few idle swings when the air in the pump is not compressed. Heating occurs as a result of the so-called adiabatic compression - when the heat does not have time to dissipate in the surrounding space. This means that when a mixture of nitrogen oxides is compressed, it must also heat up. And when heated, the equilibrium in this mixture shifts strongly towards the dioxide.

How hot will the mixture be when compressed? In the case of air compression in the pump, the heating can be easily calculated using the adiabatic equation for ideal gas: TVγ–1 = const, where T is the gas temperature (in kelvins), V is its volume, γ = C p / C v is the ratio of the heat capacity of a gas at constant pressure to the heat capacity at constant volume. For monatomic (noble) gases, γ = 1.66, for diatomic (air also belongs to them) γ = 1.40, for triatomic (for example, for NO2) γ = 1.30, etc. The adiabatic equation for air, compressible from volume 1 to volume 2 can be rewritten as T 2/ T 1 = (V 1/ V 2)γ–1. If the piston is sharply pushed to the middle of the pump, when the volume of air in it is halved, then for the ratio of temperatures before and after compression we obtain the equation T 2/ T 1 = = 20.4 = 1.31. And if T 1 \u003d 293 K (20 ° C), then T 2 = 294 K (111 ° C)!

It is impossible to directly apply the ideal gas equation to calculate the state of a mixture of nitrogen oxides immediately after compression, since in this process not only the volume, pressure and temperature change, but also the number of moles (NO2 N2O4 ratio) during the chemical reaction. The problem can be solved only by numerical integration of the differential equation, which takes into account that the work performed at each moment by the moving piston is spent, on the one hand, on heating the mixture, and on the other hand, on the dissociation of the dimer. It is assumed that the dissociation energy of N2O4, the heat capacities of both gases, the value of γ for them, and the dependence of the equilibrium position on temperature are known (all these are tabular data). The calculation shows that if the initial mixture of gases at atmospheric pressure and room temperature is quickly compressed to half the volume, then the mixture will heat up by only 13 °C. If you compress the mixture to a threefold decrease in volume, the temperature will increase by 21 ° C. And even a slight heating of the mixture strongly shifts the equilibrium position towards the dissociation of N2O4.

And then there is just a slow cooling of the gas mixture, which causes the same slow shift of the equilibrium towards N2O4 and a weakening of the color, which is observed in the experiment. The cooling rate depends on the material of the syringe walls, their thickness and other conditions of heat exchange with the surrounding air, such as drafts in the room. It is significant that with a gradual shift of the equilibrium to the right, towards N2O4, dimerization of NO2 molecules occurs with the release of heat, which reduces the rate of cooling of the mixture (similar to the freezing of water in large reservoirs at the beginning of winter does not allow the air temperature to drop rapidly).

Why did none of the experimenters feel the heating of the syringe when they pushed the plunger in? The answer is very simple. The heat capacities of the gas mixture and glass (per unit mass) do not differ very much. But the mass of the glass piston is tens and sometimes hundreds of times higher than the mass of the gas. Therefore, even if all the heat of the cooling gas mixture is transferred to the walls of the syringe, these walls will heat up by only a fraction of a degree.

The considered system with an equilibrium between two nitrogen oxides is also of practical importance. At low pressure, the mixture of NO2 and N2O4 liquefies easily. This makes it possible to use it as an effective coolant, despite its high chemical activity and corrosive effect on equipment. Unlike water, which, when receiving thermal energy, for example, from a nuclear reactor, becomes very hot and can even evaporate, the transfer of heat to a mixture of nitrogen oxides leads mainly not to its heating, but to a chemical reaction - breaking the N–N bond in the molecule N2O4. Indeed, breaking the N–N bond in one mole of a substance (92 g) without heating it requires 57.4 kJ of energy. If such energy is transferred to 92 g of water at a temperature of 20 ° C, then 30.8 kJ will go to heat the water to a boil, and the remaining 26.6 kJ will lead to the evaporation of about 11 g of water! In the case of nitrogen oxides, the mixture does not heat up much, in the colder places of the installation the circulating mixture cools slightly, the equilibrium shifts towards N2O4, and the mixture is again ready to take heat.

Electrochemical activity series of metals (voltage range, a range of standard electrode potentials) - the sequence in which the metals are arranged in order of increasing their standard electrochemical potentials φ 0 corresponding to the metal cation reduction half-reaction Me n+ : Me n+ + nē → Me

A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.

Story

The sequence of metals in the order of change in their chemical activity in in general terms already known to the alchemists. The processes of mutual displacement of metals from solutions and their surface precipitation (for example, the displacement of silver and copper from solutions of their salts by iron) were considered as a manifestation of the transmutation of elements.

Later alchemists came close to understanding the chemical side of the mutual precipitation of metals from their solutions. So, Angelus Sala in his Anatomia Vitrioli (1613) came to the conclusion that the products chemical reactions consist of the same "components" that were contained in the original substances. Subsequently, Robert Boyle proposed a hypothesis about the reasons why one metal displaces another from solution, based on corpuscular representations.

In the era of the formation of classical chemistry, the ability of elements to displace each other from compounds became an important aspect of understanding reactivity. J. Berzelius, on the basis of the electrochemical theory of affinity, built a classification of elements, dividing them into "metalloids" (now the term "non-metals" is used) and "metals" and putting hydrogen between them.

The sequence of metals according to their ability to displace each other, long known to chemists, was especially thoroughly and comprehensively studied and supplemented by N. N. Beketov in the 1860s and subsequent years. Already in 1859, he made a report in Paris on the topic "Research on the phenomena of the displacement of some elements by others." In this work, Beketov included a number of generalizations about the relationship between the mutual displacement of elements and their atomic weight, linking these processes with " the original chemical properties of the elements - what is called chemical affinity» . Beketov's discovery of the displacement of metals from solutions of their salts by hydrogen under pressure and the study of the reducing activity of aluminum, magnesium and zinc at high temperatures (metallothermy) allowed him to put forward a hypothesis about the relationship between the ability of some elements to displace others from compounds with their density: lighter simple substances capable of displacing heavier ones (therefore, this series is often also called Beketov displacement series, or simply Beketov series).

Without denying the significant merits of Beketov in the development of modern ideas about the number of activity of metals, it should be considered erroneous that prevails in the domestic popular and educational literature the idea of ​​him as the sole creator of this series. Numerous experimental data obtained at the end of the 19th century disproved Beketov's hypothesis. Thus, William Odling described many cases of "activity reversal". For example, copper displaces tin from a concentrated acidified solution of SnCl 2 and lead from an acidic solution of PbCl 2; it is also capable of dissolving in concentrated hydrochloric acid with the release of hydrogen. Copper, tin and lead are in the row to the right of cadmium, however, they can displace it from a boiling slightly acidified CdCl 2 solution.

The rapid development of theoretical and experimental physical chemistry pointed to another reason for the differences in the chemical activity of metals. With the development of modern concepts of electrochemistry (mainly in the works of Walter Nernst), it became clear that this sequence corresponds to a "series of voltages" - the arrangement of metals according to the value of standard electrode potentials. Thus, instead of a qualitative characteristic - the “tendency” of a metal and its ion to certain reactions - Nerst introduced an exact quantitative value characterizing the ability of each metal to pass into solution in the form of ions, and also to be reduced from ions to metal on the electrode, and the corresponding series was named a number of standard electrode potentials.

Theoretical basis

The values ​​of electrochemical potentials are a function of many variables and therefore show a complex dependence on the position of metals in the periodic system. Thus, the oxidation potential of cations increases with an increase in the atomization energy of a metal, with an increase in the total ionization potential of its atoms, and with a decrease in the hydration energy of its cations.

In the most general form, it is clear that metals at the beginning of periods are characterized by low values ​​of electrochemical potentials and occupy places on the left side of the voltage series. At the same time, the alternation of alkali and alkaline earth metals reflects the phenomenon of diagonal similarity. Metals located closer to the middle of the periods are characterized by large potential values ​​and occupy places in the right half of the series. A consistent increase in the electrochemical potential (from -3.395 V for a pair of Eu 2+ /Eu [ ] to +1.691 V for the Au + /Au pair) reflects a decrease in the reducing activity of metals (the ability to donate electrons) and an increase in the oxidizing ability of their cations (the ability to attach electrons). Thus, the strongest reducing agent is europium metal, and the strongest oxidizing agent is gold cations Au+.

Hydrogen is traditionally included in the voltage series, since the practical measurement of the electrochemical potentials of metals is carried out using a standard hydrogen electrode.

Practical use of a range of voltages

A number of voltages are used in practice for a comparative [relative] assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for assessing cathodic and anodic processes during electrolysis:

• Metals to the left of hydrogen are stronger reducing agents than metals to the right: they displace the latter from salt solutions. For example, the interaction Zn + Cu 2+ → Zn 2+ + Cu is possible only in the forward direction.
• Metals in the row to the left of hydrogen displace hydrogen when interacting with aqueous solutions of non-oxidizing acids; the most active metals (up to and including aluminum) - and when interacting with water.
• Metals in the row to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.
• During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of metals of moderate activity is accompanied by the release of hydrogen; the most active metals (up to aluminum) cannot be isolated from aqueous solutions of salts under normal conditions.

Table of electrochemical potentials of metals

Metal	Cation	φ 0 , V	Reactivity	Electrolysis (at the cathode):
	Li +	-3,0401	reacts with water	hydrogen is released
	Cs +	-3,026
	Rb+	-2,98
	K+	-2,931
	F+	-2,92
	Ra2+	-2,912
	Ba 2+	-2,905
	Sr2+	-2,899
	Ca2+	-2,868
	EU 2+	-2,812
	Na+	-2,71
	Sm 2+	-2,68
	Md2+	-2,40	reacts with aqueous solutions of acids
	La 3+	-2,379
	Y 3+	-2,372
	Mg2+	-2,372
	Ce 3+	-2,336
	Pr 3+	-2,353
	Nd 3+	-2,323
	Er 3+	-2,331
	Ho 3+	-2,33
	Tm3+	-2,319
	Sm 3+	-2,304
	Pm 3+	-2,30
	Fm 2+	-2,30
	Dy 3+	-2,295
	Lu 3+	-2,28
	Tb 3+	-2,28
	Gd 3+	-2,279
	Es 2+	-2,23
	AC 3+	-2,20
	Dy 2+	-2,2
	Pm 2+	-2,2
	cf2+	-2,12
	Sc 3+	-2,077
	Am 3+	-2,048
	cm 3+	-2,04
	Pu3+	-2,031
	Er 2+	-2,0
	Pr 2+	-2,0
	EU 3+	-1,991
	Lr 3+	-1,96
	cf 3+	-1,94
	Es 3+	-1,91
	Th4+	-1,899
	Fm 3+	-1,89
	Np 3+	-1,856
	Be 2+	-1,847
	U 3+	-1,798
	Al 3+	-1,700
	Md 3+	-1,65
	Ti 2+	-1,63		competing reactions: both hydrogen evolution and pure metal evolution
	hf 4+	-1,55
	Zr4+	-1,53
	Pa 3+	-1,34
	Ti 3+	-1,208
	Yb 3+	-1,205
	no 3+	-1,20
	Ti 4+	-1,19
	Mn2+	-1,185
	V2+	-1,175
	Nb 3+	-1,1
	Nb 5+	-0,96
	V 3+	-0,87
	Cr2+	-0,852
	Zn2+	-0,763
	Cr3+	-0,74
	Ga3+	-0,560